The obvious way to turn a gas into a liquid is to cool it to a temperature below its
boiling point. There is another way of condensing a gas to form a liquid, however, which
involves raising the pressure on the gas. Liquids boil at the temperature at which the
vapor pressure is equal to the pressure on the liquid from its surroundings. Raising the
pressure on a gas therefore effectively increases the boiling point of the liquid.

Suppose that we have water vapor (or steam) in a closed container at 120oC
and 1 atm. Since the temperature of the system is above the normal boiling point of water,
there is no reason for the steam to condense to form a liquid. Nothing happens as we
slowly compress the container thereby
raising the pressure on the gas until the
pressure reaches 2 atm. At this point, the system is at the boiling point of water, and
some of the gas will condense to form a liquid. As soon as the pressure on the gas exceeds
2 atm, the vapor pressure of water at 120oC is no longer large enough for the
liquid to boil. The gas therefore condenses to form a liquid, as shown in the figure
below.

In theory, we should be able to predict the pressure at which a gas condenses at a
given temperature by consultinga plot of vapor pressure vs. temperature .
In practice, every compound has a critical temperature (Tc).
If the temperature of the gas is above the critical temperature, the gas can't be
condensed, regardless of the pressure applied.

The existence of a critical temperature was discovered by Thomas Andrews in 1869 while
studying the effect of temperature and pressure on the behavior of carbon dioxide. Andrews
found that he could condense CO2 gas into a liquid by raising the pressure on
the gas, as long as he kept the temperature below 31.0oC. At 31.0oC,
for example, it takes a pressure of 72.85 atm to liquify CO2 gas. Andrews found
that it was impossible to turn CO2 into a liquid above this temperature, no
matter how much pressure was applied.

Gases can't be liquified at temperatures above the critical temperature because at this
point the properties of gases and liquids become the same, and there is no basis on which
to distinguish between gases and liquids. The vapor pressure of a liquid at the critical
temperature is called the critical pressure (Pc). The vapor
pressure of a liquid never gets larger than this critical pressure.

The critical temperatures, critical pressures, and boiling points of a number of gases
are given in the table below. There is an obvious correlation between the critical
temperature and boiling point of these gases. These properties are related because they
are both indirect measures of the force of attraction between particles in the gas phase.

Critical Temperatures, Critical Pressures and Boiling Points of
Common Gases

Gas

Tc(oC)

Pc (atm)

BP (oC)

He

-267.96

2.261

-268.94

H2

-240.17

12.77

-252.76

Ne

-228.71

26.86

-246.1

N2

-146.89

33.54

-195.81

CO

-140.23

34.53

-191.49

Ar

-122.44

48.00

-185.87

O2

-118.38

50.14

-182.96

CH4

-82.60

45.44

-161.49

CO2

31.04

72.85

-78.44

NH3

132.4

111.3

-33.42

Cl2

144.0

78.1

-34.03

The experimental values of the critical temperature and pressure of a substance can be
used to calculate the a and b constants in the van der Waals
equation.

There is a force of attraction between molecules in liquids, and liquids can flow until
they take on the shape that maximizes this force of attraction. Below the surface of the
liquid, the force of cohesion (literally, "sticking together")
between molecules is the same in all directions, as shown in the figure below. Molecules
on the surface of the liquid, however, feel a net force of attraction that pulls them back
into the body of the liquid. As a result, the liquid tries to take on the shape that has
the smallest possible surface area the
shape of a sphere. The magnitude of the force that controls the shape of the liquid is
called the surface tension. The stronger the bonds between the molecules
in the liquid, the larger the surface tension.

There is also a force of adhesion (literally, "sticking")
between a liquid and the walls of the container. When the force of adhesion is more than
half as large as the force of cohesion between the liquid molecules, the liquid is said to
"wet" the solid. A good example of this phenomenon is the wetting of paper by
water. The force of adhesion between paper and water combined with the force of cohesion
between water molecules explains why sheets of wet paper stick together.

Water wets glass because of the force of adhesion that results from interactions
between the positive ends of the polar water molecules and the negatively charged oxygen
atoms in glass. As a result, water forms a meniscus that curves upward in
a small-diameter glass tube, as shown in the figure below. (The term meniscus
comes from the Greek word for "moon" and is used to describe anything that has a
crescent shape.) The meniscus that water forms in a buret results from a balance between
the force of adhesion pulling up on the column of water to wet the walls of the glass tube
and the force of gravity pulling down on the liquid.

Water climbs the walls of a small-diameter tube to form a meniscus that curves upward,
whereas mercury forms a meniscus that curves downward.

The force of adhesion between water and wax is very small compared to the
force of cohesion between water molecules. As a result, rain doesn't adhere to
wax. It tends to form beads, or drops, with the smallest possible surface area, thereby
maximizing the force of cohesion between the water molecules. The same thing happens when
mercury is spilled on glass or poured into a narrow glass tube. The force of cohesion
between mercury atoms is so much larger than the force of adhesion between mercury and
glass that the area of contact between mercury and glass is kept to a minimum, with the
net result being the meniscus shown in the above figure.

Viscosity is a measure of the resistance to flow. Motor oils are more
viscous than gasoline, for example, and the maple syrup used on pancakes is more viscous
than the vegetable oils used in salad dressings.

Viscosity is measured by determining the rate at which a liquid or gas flows through a
small-diameter glass tube. In 1844 Jean Louis Marie Poiseuille showed that the volume of
fluid (V) that flows down a small-diameter capillary tube per unit of time (t)
is proportional to the radius of the rube (r), the pressure pushing the fluid
down the tube (P), the length of the tube (l), and the viscosity of the
fluid ().

Viscosity is reported in units called poise (pronounced "pwahz").
The viscosity of water at room temperature is roughly 1 centipoise, or 1 cP. Gasoline has
a viscosity between 0.4 and 0.5 cP; the viscosity of air is 0.018 cP.

Because the molecules closest to the walls of a small-diameter tube adhere to the
glass, viscosity measures the rate at which molecules in the middle of the stream of
liquid or gas flow past this outer layer of more or less stationary molecules. Viscosity
therefore depends on any factor that can influence the ease with which molecules slip past
each other. Liquids tend to become more viscous as the molecules become larger, or as the
amount of intermolecular bonding increases. They become less viscous as the temperature
increases. The viscosity of water, for example, decreases from 1.77 cP at 0oC
to 0.28 cP at 100oC.

We are so familiar with the properties of water that it is difficult to appreciate the
extent to which its behavior is unusual.

Most solids expand when they melt. Water expands when it freezes.

Most solids are more dense than the corresponding liquids. Ice (0.917 g/cm3)
is not as dense as water.

Water has a melting point at least 100oC higher than expected on the basis of
the melting points of H2S, H2Se, and H2Te.

Water has a boiling point almost 200oC higher than expected from the boiling
points of H2S, H2Se, and H2Te.

Water has the largest surface tension of any common liquid except liquid mercury.

Water has an unusually large viscosity.

Water is an excellent solvent. It can dissolve compounds, such as NaCl, that are
insoluble or only slightly soluble in other liquids.

Water has an unusually high heat capacity. It takes more heat to raise the temperature
of 1 gram of water by 1oC than any other liquid.

These anomalous properties all result from the strong intermolecular bonds in water.
Water is best described as a polar
molecule in which there is a partial separation of charge to give positive and
negative poles. The force of attraction between a positively charged hydrogen atom on one
water molecule and the negatively charged oxygen atom on another gives rise to an
intermolecular bond, as shown in the figure below. This dipole-dipole interaction between
water molecules is known as a hydrogen bond.

Hydrogen bonds are separated from other examples of van der Waals forces because they
are unusually strong: 10-12 kJ/mol. The hydrogen bonds in water are particularly important
because of the dominant role that water plays in the chemistry of living systems. Hydrogen
bonds are not limited to water, however.

Hydrogen-bond donors include substances that contain relatively polar H-X
bonds, such as NH3, H2O, and HF. Hydrogen-bond acceptors include
substances that have nonbonding pairs of valence electrons. The H-X bond must be
polar to create the partial positive charge on the hydrogen atom that allows dipole-dipole
interactions to exist. As the X atom in the H-X bond becomes less
electronegative, hydrogen bonding between molecules becomes less important. Hydrogen
bonding in HF, for example, is much stronger than in either H2O or HCl.

The hydrogen bonds between water molecules in ice produce the open structure shown in
the figure below. When ice melts, some of these bonds are broken, and this structure
collapses to form a liquid that is about 10% denser. This unusual property of water has
several important consequences. The expansion of water when it freezes is responsible for
the cracking of concrete, which forms potholes in streets and highways. But it also means
that ice floats on top of rivers and streams. The ice that forms each winter therefore has
a chance to melt during the summer.

The structure of ice. Note that the hydrogen atoms are closer to one of the oxygen
atoms than the other in each of the hydrogen bonds.

The figure below shows another consequence of the strength of the hydrogen bonds in
water. There is a steady increase in boiling point in the series CH4, GeH4,
SiH4, and SnH4. The boiling points of H2O and HF,
however, are anomalously large because of the strong hydrogen bonds between molecules in
these liquids. If this doesn't seem important, try to imagine what life would be like if
water boiled at -80oC.

The surface tension and viscosity of water are also related to the strength of the
hydrogen bonds between water molecules. The surface tension of water is responsible for
the capillary action that brings water up through the root systems of plants. It is also
responsible for the efficiency with which the wax that coats the surface of leaves can
protect plants from excessive loss of water through evaporation.

The unusually large heat capacity of water is also related to the strength of the
hydrogen bonds between water molecules. Anything that increases the motion of water
molecules, and therefore the temperature of water, must interfere with the hydrogen bonds
between these molecules. The fact that it takes so much energy to disrupt these bonds
means that water can store enormous amounts of thermal energy. Although the water in lakes
and rivers gets warmer in the summer and cooler in the winter, the large heat capacity of
water limits the range of temperatures that would otherwise threaten the life that
flourishes in this environment. The heat capacity of water is also responsible for the
ocean's ability to act as a thermal reservoir that moderates the swings in temperature
that occur from winter to summer.